Chapter 8 Periodicity Flashcards
(19 cards)
how are the elements arranged in the periodic table?
by the proton number NOT the mass number
- elements in the same group will have similar properties
- elements in the same period have the same number of electron shells
what is the trend of atomic radii in period 3? why?
as we go across period 3, the atomic radius decreases.
this is due to an increasing nuclear charge as there is an increasing number of protons but the electron shell number remains the same. therefore this pulls the electrons further in towards the nucleus
trend of atomic radius down a group?
increases down a group due to the addition of extra electron shells
what is the trend of melting point across period 3?
first 3 elements are metals
There is a general increase in melting points as metal ions have an increasing positive charge thus increasing the number of delocalised electrons and smaller ionic radius -> stronger bond
why does silicon have a higher melting point than aluminium?
has a giant covalent structure, therefore many strong covalent bonds hold the silicon atoms together. large amount of energy is needed to break these
why does phosphorus have a lower melting point than silicon?
has the formula P4
will have a lower melting point due to a weaker simple molecular structure. therefore melting point is determined by the presence of the weaker van der waals forces
why does the melting point rise from phosphorus to sulphur?
has the formula S8
will have a higher melting point as it is a larger simple molecular structure, so larger van der waals forces
why does chlorine have a lower melting point than sulphur?
formula Cl2
smaller simple molecular structure so smaller van der waals.
why does argon have a low melting point?
exists as an individual atom therefore has smaller van der waals than chlorine, so lower melting point
what is ionisation energy?
the minimum amount of energy required to remove 1 mole of electrons from 1 mole of atoms in the gaseous state
what factors impact ionisation energies?
shielding, atomic size and nuclear charge
how does shielding affect ionisation energies?
the more electron shells between the positive nucleus and the negative electron that is being removed, the less energy is required as there is a weaker attraction
how does atomic size affect ionisation energies?
the bigger the atom, the further away the outer electron from the nucleus, therefore attractive force between the two reduces, thus easier to remove = less energy
how does nuclear charge affect ionisation energies?
the more protons in the nucleus, the bigger the attraction between nucleus and outer electrons. therefore more energy required to remove the electron
what is successive ionisation?
the remove of more than 1 electron from the same atom
what is the trend of successive ionisation across a period?
general increase in energy as removing an electron from an increasingly more positive ion
there will be jumps in energy as removing electrons from shell closer to the nucleus requires more energy
what is the trend in ionisation energy down a group?
decreases
- the atomic radius increases, therefore outer electrons are further from the nucleus thus the attractive force is weaker.
energy required decreases
-shielding increases as there are more shells, attractive force is weaker
model provides evidence for shells
what is the trend of ionisation energy across a period?
increases
- increasing number of protons, so increase in nuclear attraction
- shielding remains the same/ similar so no impact
- more energy required
what are the exceptions for ionisation energy across a period?
evidence for sub-shells
aluminium - in a higher energy sub-shell which is slightly further from nucleus
sulphur - decrease as evidence for electron repulsion
involves trying to remove from an orbital with 2 electrons, eletrons repel each other so less energy needed unlike Phosphorus
